The Periodic Law
Electron Configuration
and Periodic Properties
Electron Configuration and Periodic Properties
Objectives:
1. Define atomic radius and ionization energy.
2. Compare the periodic trends of atomic radii
and ionization energy.
Atomic Radius
Atomic Radius – distance
from the nucleus to the
outermost electron.
• can be measured as onehalf the distance between
nuclei of identical atoms
bonded together
Atomic Radius
Atomic Radius
What are the trends on the periodic table?
1.Atomic radius increases going down a
group on the P.T.
2.Atomic radius decreases moving left to
right across a period.
Atomic Radius
How can these trends be explained?
Group Trend
1. Outer electrons occupy higher energy levels;
electrons are farther from the nucleus
Period Trend
2. Increasing nuclear charge causes atoms to be
smaller (left to right); increased pull on outer
electrons due to nuclear charge.
Ionization Energy
Ionization energy – energy required to remove one
electron from a neutral atom (first ionization
energy)
X + energy  X+ + e• Measured using isolated atoms in the gas phase.
Ionization Energy
What are the trends on the periodic table?
1. I.E. increases going left to right across
a period.
2. I.E increases from bottom to the top of
a group
Ionization Energy
How can these trends be explained?
Period Trend
I.E. increases left to right as a result of increasing nuclear
charge; a larger nuclear charge results in stronger
attraction for electrons
Group Trend
I.E decreases going down a group due to the fact that
electrons are in higher energy levels and farther from the
nucleus; also due to the shielding effects of inner
electrons.
Successive Ionization Energies
Electron Affinity
• The energy change that occurs when a neutral
atom acquires an electron.
X + e-  X- + energy (- energy change)
or
X + e- + energy  X- (+ energy change)
Electron Affinity Trends
• Halogens have the highest electron affinities.
Fluorine - 1s22s22p5
• Electron affinity increases moving to the right on the
periodic table. (not always & not including noble
gases)
Carbon – 1s22s22p2
Nitrogen - 1s22s22p3
• Electron affinity increases moving up a group on the
p.t. This trend is not as clear as it is for ionization
energy. E.A is affected by nuclear charge and atomic
radius
Ionic Radii
•When atoms become
cations, they become
smaller.
•When atoms become
anions, they become larger.
Valence Electrons
Valence Electrons – the electrons located in s and p
orbitals for main group elements
• Atoms gain, lose, or share electrons to in order to
have a complete set of s and p electrons.
N
P
1s22s22p3
1s22s22p63s23p3
Electronegativity
Electronegativity – the ability of an atom to attract
electrons in a chemical bond.
Electronegativity Trends
• N, O, and the halogens have the highest
electronegativities
• Electronegativity increases going left to right across
the p.t. up until the halogens
• Electronegativity increases going up groups on the
p.t.
Periodic Properties
1. Which element (Cs, Hf, Au) has the smallest
atomic radius?
Au
2. Arrange the following elements in order of
decreasing electron affinity: C, O, Li, Na, Rb, F
F, O, C, Li, Na, Rb
Periodic Properties
3. Arrange the following elements in order of
decreasing first ionization energy: Li, O, C, K,
Ne, F
Ne, F, O, C, Li, K
4. Which element in #3 would have the highest
second ionization energy and why?
Li
Periodic Properties
5. Which element is the most electronegative
among C, N,O, and S?
O
6. Which ion has the smallest radius, K+ or Ca2+?
Ca2+